As used herein, the term “supersorbent” shall mean a sorbent as taught in U.S. Pat. No. 5,779,464 entitled “Calcium Carbonate Sorbent and Methods of Making and Using Same”, the teachings of which are hereby incorporated by reference.
As used herein, the term “microporous” shall mean a pore size distribution of less than 5 nanometers. As used herein, the term “mesoporous” shall mean a pore size distribution of from about 5 nanometers to about 20 nanometers.
Atmospheric CO2 concentration has been increasing steadily since the industrial revolution. It has been widely accepted that the while the CO2 concentration was about 280 ppm before the industrial revolution, it has increased from 315 ppmv in 1959 to 370 ppmv in 2001 [Keeling, C. D. and T. P. Whorf. 2002. Atmospheric CO2 records from sites in the SIO air sampling network. In Trends: A Compendium of Data on Global Change. Carbon Dioxide Information Analysis Center, Oak Ridge National Laboratory, U.S. Department of Energy, Oak Ridge, Tenn., U.S.A. This data is also available from http://cdiac.esd.ornl.gov/ftp/maunaloa-co2/maunaloa.co2]. Rising CO2 concentrations has been reported to account for half of the greenhouse effect that causes global warming [IPCC Working Group I. IPCC Climate Change 1995—The Science of Climate Change: The Second Assessment Report of the Intergovernmental Panel on Climate Change; Houghton, J. T., Meira Filho, L. G., Callander, B. A., Harris, N., Kattenberg, A., Maskell K, Eds.; Cambridge University Press: Cambridge, U.K., 1996]. Although the anthropogenic CO2 emissions are small compared to the amount of CO2 exchanged in the natural cycles, the discrepancy between the long life of CO2 in the atmosphere (50-200 years) and the slow rate of natural CO2 sequestration processes leads to CO2 build up in the atmosphere. The IPCC (Intergovernmental Panel on Climate Change) opines that “the balance of evidence suggests a discernible human influence on the global climate.” Therefore, it is necessary to develop cost effective CO2 management schemes to curb its emission.
Many of the envisaged CO2 management schemes consist of three parts—separation, transportation and sequestration of CO2 [FETC Carbon Sequestration R&D Program Plan: FY 1999-2000. National Energy Technology Laboratory, Department of Energy, Washington, D.C., 1999]. The cost of separation and compression of CO2 to 110 bar (for transportation of CO2 in liquid state) is estimated at $30-50 per ton CO2, and transportation and sequestration would cost about $1-3 per ton per 100 km and $1-3 per ton of CO2, respectively [Wallace, D. Capture and Storage of CO2. What Needs To Be Done. Presented at the 6th Conference of the Parties, COP 6, to the United Nations Framework Convention on Climate Change; The Hague, The Netherlands, Nov. 13-24, 2000; www.iea.org/envissu/index.htm]. The capture of CO2 imposes severe energy penalties thereby reducing the net electricity output by as much as 13-37% [Herzog, H.; Drake, E.; Adams, E. CO2 Capture, Reuse, and Storage Technologies for Mitigating Global Climate Change. A White Paper; Final Report No. DE-AF22-96PC01257, January 1997]. The dominating costs associated with the current CO2 separation technologies necessitate development of economical alternatives.
Historically, CO2 separation was motivated by enhanced oil recovery [Kaplan, L. J. Cost-Saving Processes Recovers CO2 from Power-Plant Flue gas. Chem. Eng. 1982, 89 (24), 30-31; Pauley, C. P.; Smiskey, P. L.; Haigh, S. N-ReN Recovers CO2 from Flue Gas Economically. Oil Gas J. 1984, 82(20), 87-92]. Currently, industrial processes such as limestone calcination, synthesis of ammonia and hydrogen production require CO2 separation. Absorption processes employ physical and chemical solvents such as Selexol and Rectisol, MEA and KS-2 [Reimer, P.; Audus, H.; Smith, A. Carbon Dioxide Capture from Power Stations. IEA Greenhouse R&D Programme, www.ieagreen.org.uk, 2001. ISBN 1 898373 15 9; Blauwhoff, P. M. M.; Versteeg, G. F.; van Swaaij, W. P. M. A study on the reaction between CO2 and alkanoamines in aqueous solution. Chem. Eng. Sci. 1984, 39(2), 207-225. Mimura, T.; Simayoshi, H.; Suda, T.; Iijima, M.; Mitsuake, S. Development of Energy Saving Technology for Flue Gas Carbon Dioxide Recovery by Chemical Absorption Method and Steam System in Power Plant. Energy Convers. Mgmt. 1997, 38, Suppl. P. S57-S62]. Adsorption systems capture CO2 on a bed of adsorbent materials such as molecular sieves and activated carbon [Kikkinides, E. S.; Yang, R. T.; Cho, S. H. Concentration and Recovery of CO2 from flue gas by pressure swing adsorption. Ind. Eng. Chem. Res. 1993, 32, 2714-2720]. CO2 can also be separated from the other gases by condensing it out at cryogenic temperatures. Polymers, metals such as palladium, and molecular sieves are being evaluated for membrane based separation processes [Reimer, P.; Audus, H.; Smith, A. Carbon Dioxide Capture from Power Stations. IEA Greenhouse R&D Programme, www.ieagreen.org.uk, 2001. ISBN 1 898373 15 9].
Reaction based processes, as promulgated in this work, can be applied to separate CO2 from gas mixtures. This process is based on a heterogeneous gas-solid non-catalytic carbonation reaction where gaseous CO2 reacts with solid metal oxide (represented by MO) to yield the metal carbonate (MCO3). The reaction can be represented by:MO+CO2→MCO3  (1)Once the metal oxide has reached its ultimate conversion, it can be thermally regenerated to the metal oxide and CO2 by the calcination of the metal carbonate product. The calcination reaction can be represented by:MCO3→MO+CO2  (2)As an example of the above-mentioned scheme, FIG. 1 shows the variation in the free energy of the carbonation reaction as a function of temperature for calcium oxide. From the figure, we can see that the carbonation reaction is thermodynamically favored with a decrease in temperature (Gibbs free energy declines with a decrease in temperature). However, at lower temperatures, the carbonation reaction is kinetically slow. In fact, it takes geological time scales for the formation of CaCO3 by the reaction between CaO and atmospheric CO2 (at 280-360 ppm) at ambient temperatures. It should also be noted that the carbonation reaction would be favored as long as the free energy is negative. This creates an upper bound of 890° C. for carbonation to occur under a CO2 partial pressure of 1 atm. The equilibrium temperature for this reaction is a function of the partial pressure of CO2. A reaction based CO2 separation process offers many advantages. Under ideal conditions, MEA captures 60 g CO2/kg, silica gel adsorbs 13.2 g CO2/kg and activated carbon adsorbs 88 g CO2/kg. The sorption capacity of some metal oxides (such as the modified CaO, presented in this study) is about 700 g CO2/kg of CaO. This is about an order of magnitude higher than the capacity of adsorbents/solvents used in other CO2 separation processes and would significantly reduce the size of the reactors and the material handling associated with CO2 separation.
Numerous metal oxides exhibit the carbonation and calcination reaction. The calcination temperature of a few metal carbonates (CaCO3 ˜750° C., MgCO3 ˜385° C., ZnCO3 ˜340° C., PbCO3 ˜350° C., CuCO3 ˜225-290° C. and MnCO3 ˜440° C.) makes them viable candidates for this process. Apart from CaO, gas-solid carbonation of other metal oxides has not been widely studied. The carbonation of ZnO to ZnCO3 at 8-13° C. was low when exposed to CO2 and H2O for over 100 days (Sawada, Y.; Murakami, M.; Nishide, T. Thermal analysis of basic zinc carbonate. Part 1. Carbonation process of zinc oxide powders at 8 and 13° C. Thermochim. Acta. 1996, 273, 95-102.). MnCO3 undergoes a more complex thermal degradation phenomena. MnCO3 first decomposes to MnO2 at 300° C., which in turn changes to Mn2O3 at 440° C. At higher temperatures (˜900° C.), the final thermal decomposition product was identified as Mn3O4 (Shaheen, W. M.; Selim, M. M. Effect of thermal treatment on physicochemical properties of pure and mixed manganese carbonate and basic copper carbonate. Thermochim. Acta. 1998, 322(2), 117-128.). Different oxides of manganese provide the flexibility of exploiting the carbonation/calcination reaction over a wider temperature range. Aqueous phase MgO carbonation has been studied for its suitability for mineral-based CO2 sequestration (Fernandez, A. I.; Chimenos, J. M.; Segarra, M.; Fernandez, M. A.; Espiell, F. Kinetic study of carbonation of MgO slurries. Hydrometallurgy. 1999, 53, 155-167). The carbonation extent of Mg(OH)2 was about 10% between 387-400° C. and 6% formation between 475-500° C. (Butt, D. P.; Lackner, K. S.; Wendt, C. H.; Conzone, S. D.; Kung, H.; Lu, Y-C.; Bremser, J. K. Kinetics of Thermal Dehydroxylation and Carbonation of Magnesium Hydroxide. J. Am. Ceram. Soc. 1996, 79(7), 1892-1898). They attributed the low conversions to the formation of a non-porous carbonate product layer. This layer hinders the inward diffusion of CO2 and the outward diffusion of H2O (a product of the carbonation reaction) leading to low conversions. The carbonation of PbO was studied as a part of the chemical heat pump process (Kato, Y.; Saku, D.; Harada, N.; Yoshizawa, Y. Utilization of High Temperature Heat from Nuclear Reactor using Inorganic Chemical Heat Pump. Progress in Nuclear Energy. 1998, 32(3-4), 563-570. & Kato, Y.; Harada, N.; Yoshizawa, Y. Kinetic feasibility of a chemical heat pump for heat utilization from high temperature processes. Applied Thermal Engineering. 1999, 19, 239-254). They reported 30% conversion in an hour under 100% CO2 atmosphere at 300° C. Furthermore, they found the reactivity of PbO to drop with the number of carbonation-calcination cycles.
Carbonation of calcium oxide has been widely studied. Related applications of the CaO carbonation and calcination include the storage of energy (Barker, R. The Reversibility of the Reaction CaCO3=CaO+CO2. J. Appl. Chem. Biotechnol. 1973, 23, 733-742) and the zero emission coal alliance process, consisting of hydrogasification of coal fueled by the heat of the carbonation reaction (Tinkler, M. J.; Cheh, C. Towards a Coal-capable Solid Oxide Fuel Cell System. Proceedings of the 26th International Technical Conference on Coal Utilization and Fuel Systems; Clearwater, Fla., Mar. 5-8, 2001; pp 569-570). The gas-solid CaO—CO2 reaction proceeds through two rate-controlling regimes. The first regime involves a rapid, heterogeneous chemical reaction. In the second regime, the reaction slows down due to the formation of an impervious layer of CaCO3. This product layer prevents the exposure of unreacted CaO in the particle core to CO2 for further carbonation. The kinetics of the second regime is governed by the diffusion of ions through the CaCO3 product layer. The activation energy was estimated to be 21 kcal/mol below 688 K and 43 kcal/mol above it for the product layer diffusion, based on the counter migration of CO32− and O2− ions through the product layer (Bhatia, S. K.; and Perlmutter, D. D. Effect of the product layer on the kinetics of the CO2-Lime Reaction. AlChE J. 1983, 29(1), 79-86).
The extent of the carbonation reaction reported in many studies has also shown considerable variation. Stoichiometrically, 56 g of CaO should react with 44 g of CO2 to form 100 g of CaCO3. This translates to about 78.6-wt % capacity for CaO. However, the structural limitations prevent the attainment of theoretical conversion. The extent of carbonation was only 23-wt % in 30 minutes at 600° C. (Dedman, A. J.; Owen, A. J. Calcium Cyanamide Synthesis, Part 4.—The reaction CaO+CO2=CaCO3. Trans. Faraday Soc. 1962, 58, 2027-2035). A higher surface area CaO sorbent provided 55-wt % CO2 sorption (Bhatia, S. K.; and Perlmutter, D. D. Effect of the product layer on the kinetics of the CO2-Lime Reaction. AlChE J. 1983, 29(1), 79-86). 64-wt % CO2 sorption was achieved at 1050° C. temperature and 11.74 atm CO2 pressure in 32 hours (Mess, D.; Sarofim, A. F.; Longwell, J. P. Product Layer Diffusion during the Reaction of Calcium Oxide with Carbon Dioxide. Energy and Fuels. 1999, 13, 999-1005). However, the extent of carbonation at lower temperature/pressure conditions that are more characteristic of CO2 containing gaseous mixtures is absent in their work. The limitation in total conversion stems essentially from the nature of the initial pore size distribution of the CaO sorbent. Microporous sorbents (pore size <2 nm) are very susceptible to pore blockage and plugging due to the formation of higher molar volume product (molar volume of CaO: 17 cm3/mol; molar volume of CaCO3: 37 cm3/mol). CaO sorbents obtained from naturally occurring precursors are usually microporous in nature. At the end of the kinetically controlled regime, diffusion processes through the product layer control the reaction rate. Similar structural limitations have prevented calcium-based sorbents from attaining theoretical conversion for the sulfation reaction between CaO and sulfur dioxide (SO2) as well (Wei, S.-H.; Mahuli, S. K.; Agnihotri, R.; Fan, L.-S. High Surface Area Calcium Carbonate: Pore Structural Properties and Sulfation Characteristics. Ind. Eng. Chem. Res. 1997, 36(6), 2141-2148). They suggested that a mesoporous structure, which maximizes porosity in the 5-20 nm pore size range, would be less susceptible to pore pluggage. This structure would also be able to provide sufficient surface area to ensure rapid kinetics. Their modified precipitation technique resulted in a mesoporous CaCO3 structure that also had a high BET surface area determined by nitrogen (60 m2/g). A similar approach could also enhance the reactivity of CaO sorbents towards the carbonation reaction, which is the focus of this study.
Lastly, it is important that the CaO sorbents maintain their reactivity over many carbonation and calcination cycles. The conversion of CaO dropped from about 73% in the first carbonation cycle to 43% at the end of the 5th cycle at 866° C. (Barker, R. The Reversibility of the Reaction CaCO3=CaO+CO2. J. Appl. Chem. Biotechnol. 1973, 23, 733-742 & Barker, R. The Reactivity of Calcium Oxide Towards Carbon Dioxide and its use for Energy Storage. J. Appl. Chem. Biotechnol. 1974, 24, 221-227). Barker suggested that the CaCO3 layer is about 22 nm thick and his latter work showed repeated 93% conversion over 30 cycles at 629° C. on 10 nm CaO particles. In another study, cyclical studies conducted at a carbonation temperature of 880° C. and calcination at 860° C. led to a drop in conversion from 70% in the first carbonation to 38% in the 7th carbonation step (Kato, Y.; Harada, N.; Yoshizawa, Y. Kinetic feasibility of a chemical heat pump for heat utilization from high temperature processes. Applied Thermal Engineering. 1999, 19, 239-254). The process described here leads to >95% conversion due to the application of novel mesoporous CaO sorbents for CO2 capture and maintains their reactivity over repeated cycles of carbonation and calcination.
Part I (CO2/SO2 Combined Reaction Optimization)
Introduction
Carbon dioxide (CO2) accounts for more than half of the enhanced greenhouse effect, which is responsible for global warming. ‘The atmospheric concentration of CO2 has increased from 280 ppm before the Industrial Revolution to −365 ppm today. 2’2,3 This is mainly due to the unabated emission of CO2 as a result of increasing consumption of fossil fuels such as coal, oil and natural gas. Point sources, such as electric utility plants that contribute to about one-third of all anthropogenic CO2 emissions4, are ideal candidates for implementing CO2 reduction practices due to the relatively high concentration and quantity of CO2 emitted compared to smaller, mobile sources. Coal consumption leads to high C02 emissions at these large point sources due to its dominant use in electricity generation (−52%) and higher energy specific CO2 emission due to its high carbon to hydrogen content compared to other fossil fuels (g C02BTU).s Comprehensive CO2 management scenarios involve a three-step process that includes separation, transportation and safe sequestration of CO2. Economic analysis has, shown that CO2 separation accounts for 75-85% of the overall cost associated with carbon sequestration.6 Current CO2 separation technologies based on absorption, adsorption, membrane separation, and cryogenic separation necessitate a low temperature and/or high pressure of flue gas to enhance the CO2 sorption capacity of the sorbent/solvent or the diffusion flux of CO2 through the membrane. However, flue gas is typically characterized by sub-atmospheric pressure and high temperature. Metal oxides are capable of reacting with CO2 under existing flue gas conditions, thereby reducing downstream process modifications. We have detailed elsewhere the advantages of a high temperature reactive separation process based on the carbonation and calcination reactions (CCR) of CaO to separate CO2 from flue gas.7 The key advantage offered by this process is the enhanced C02 sorption capacity (35-70 weight %) exhibited by the high reactivity CaO particles under existing flue gas conditions over multiple cycles of CCRs.
Extensive screening of metal oxides has identified CaO as a potential candidate for the CCR scheme.7 The carbonation reaction of CaO has been studied for its role in chemical heat pumps8,9, energy storage systems10, zero emission coal alliance processes”, and in the enhanced production of hydrogen from fossil fuels. ‘This reaction typically goes through a raid kinetic controlled regime, followed by a slower product-layer diffusion controlled regime.’ Naturally occurring precursors (limestone and dolomite), are unable to achieve stoichiometric conversion in any carbonation step due to the predominant microporous structure which is susceptible to pore pluggage and pore mouth closure. In contrast, mesoporosity, which dominates the pore structure of precipitated calcium carbonate (PCC), synthesized under the influence of negatively charged polyacrylate ions yields greater than 90% carbonation conversion.7,4 
For the viability of a CCR process, it is imperative that the CaO sorbent maintain high reactivity over multiple cycles. Previous studies in the literature have reported the performance of numerous CaO sorbents over multiple cycles. Abanades and co-workers summarized the CCR experimental data of previous studies on a variety of CaO sorbents differing in their physical properties. They were able to develop a single correlation between the extent of carbonation as a function of the number of CCR cycles.15,16 These sorbents experienced a similar loss in reactivity towards the carbonation reaction regardless of differences in particles size, reactor types, reaction conditions, sorbent characteristics and cycle times. They observed that the highest C02 sorption capacity retained by the sorbent was 24 wt % after 20 cycles.
Sulfur present in coal oxidizes to S02 during combustion. Calcium based sorbents are widely used for the control of S02 emissions. The two principal calcium utilization processes are low temperature wet scrubbing and high temperature furnace sorbent injection (FSI). In wet scrubbing, SO2 capture occurs through ionic reactions in the aqueous phase. In high temperature (>900° C.) FSI systems, calcium oxide precursors (dolomite, Ca(OH)2 and limestone) and their calcines reacts with SO2 to form CaSO4 via the heterogeneous non-catalytic gas solid reaction. Sulfation under these conditions has been extensively studied and simulated using various models.”, “,” PCC achieves a higher extent of sulfation (−70%) compared to naturally occurring limestone (−30%) at greater than 900° C. within a residence time of 700 milliseconds.
The flue gas generated by coal combustion typically contains 10-15% C02, 3-4% 02, 5-7% H2O, 500-3000 ppm SO2 and 150-500 ppm NOx in addition to trace quantities of HCl, arsenic, mercury, and selenium. Separation of CO2 by its absorption in monoethanolamine (MEA) is currently the most viable option for commercial scale deployment. However, MEA forms thermally stable salts with SO2 and NOx, which do not decompose under the regeneration conditions employed in the MEA process. It is necessary to lower SO2 concentration to below 10 ppm to minimize the loss of the costly solvent. Economic analysis of this process, based on a parasitic consumption of MEA of 0.5-2 g MEA/kg C02 separated, show that the cost associated with CO2 separation lies in the $33-73/ton CO2 avoided.21 A similar hurdle is posed by SO2 for a CaO based CCR process. CaO undergoes sulfation with SO2 forming CaSO4, which cannot be thermally decomposed back to CaO within the operating temperature range of the proposed CCR process (400-800° C.) as it requires greater than 1100° C. for its decomposition. Exposure of CaO to higher temperatures leads to a loss in surface area and porosity due to excessive sintering, which drastically reduces its reactivity. Eventually, the CaSO4 buildup in each cycle reduces the regenerative capacity of the CaO sorbent over subsequent cycles ultimately rendering it inactive. However, literature on the sulfation of CaO in the temperature range where CaCO3 is thermodynamically stable is scant. Sulfation of calcium species in this temperature range is crucial for the experiments covered in this paper because this study aims to investigate the effect of SO2 on the carbonation of CaO.
The simultaneous hydration, carbonation and sulfation of reagent grade 5 micron CaO particles have been previously investigated for an exposure time of 2 hours in the 170-580° C. temperature range under differential conditions.22 Low temperatures favor hydration over sulfation and carbonation. 380° C. marks the termination of hydration and the onset of carbonation and sulfation. Carbonation peaked at 520° C. whereas sulfation dominated beyond 580° C. Furthermore, the sulfur species in the form of CaSO3 peaks at 24% at 300° C. and CaSO4 is the only sulfur species above 585° C. A high extent of sulfation has also been attained by 2 mm sized macroporous (>200 run) CaO particles synthesized by a swelling technique involving water-acetic acid mixtures.23.11 The authors attributed the high sulfation extent to the increased access of SO2 to the particle surface due to the macroporosity of the sorbent. Li et al. investigated combined carbonation and sulfation reactions on commercial grade calcium hydroxide.25 They carried out these reactions at 425-650° C. by exposing the fines for 2 seconds under entrained flow conditions. The particles were then collected on a hot filter, maintained between 450-510° C. and further exposed to the gas mixture for 2 hours. Their results indicate an increasing extent of direct sulfation of the carbonated product (CaCO3) with higher residence time.2S The kinetic analysis and modeling of the reaction between SO2 and CaCO3 in the temperature range where CaCO3 is thermodynamically stable, was studied by Snow et al. and Hajaligol et al.26,27 They exploited the higher porosity of calcium oxalate derived CaCO3 to achieve about 90% sulfation in the 400-550° C. temperature range. Tullin and Ljungstrom conducted thermogravimetric studies on the simultaneous carbonation and sulfation of CaO and CaCO3 for a residence time of 10-180 minutes at 860° C. The gas mixture consisted of 30-80% CO2i 3000 ppm SO2 and 3-4% oxygen28,29 The initial increase in weight of the sorbent was predominantly due to the carbonation reaction, which occurs to a higher extent than sulfation for the given inlet gas concentration levels. Further exposure of the sorbent to the reactant gas mixture results in the direct sulfation of the CaCO3 so formed, and leads to a decrease in the overall extent of carbonation and an increase in sulfation. In other experiments, they show that although both CaO and CaCO3 have similar reactivity towards the sulfation reaction, CaSO4 formed due to the direct carbonation of CaCO3 is more porous than the CaSO4 product formed by sulfation of CaO.
These literature studies indicate that carbonation occurs at faster rate compared to sulfation at temperatures around 700° C. due to higher concentration of CO2. However, SO2 will eventually react directly with CaCO3 leading to the formation of CaSO4. It is thus imperative to obtain experimental data on combined carbonation and sulfation reactions of CaO over multiple cycles to identify the process conditions under which the extent of carbonation can be maximized in the presence of SO2. Simultaneous high temperature carbonation and sulfation experiments were performed in a Thermogravimetric Analyzer (TGA). The study demonstrates the effect of solid residence time on the overall extent of simultaneous carbonation and sulfation.
Enhanced Hydrogen Production Integrated with CO2 Separation in a Single-Stage Reactor
There has been a global push towards the development of a hydrogen economy. The main premise behind this drastic alteration in our energy usage stems from the fact that the use of hydrogen in portable and mobile applications would be the most environmentally beneficial process that leads only to the emission of water. However, the biggest issue that needs to be addressed for the success of the hydrogen-based economy involves the source of hydrogen itself. While hydrogen may be considered as the best “carrier” of energy, there is clearly no hydrogen “wells” on earth. The major processes for hydrogen production from fossil fuels consist of steam reforming of methane (SMR), coal gasification, catalytic cracking of natural gas, and partial oxidation of heavy oils. Other processes consist of water electrolysis, thermo chemical water decomposition, biological processes, etc. (Rosen and Scott, 1998; Rosen, 1996). However, water electrolysis is not a very energy efficient process.
Water gas, a mixture of CO, CO2, H2O and H2, is formed by the gasification of coal by sub-stoichiometric air and/or steam. Irrespective of the initial concentration of these four gases, the reversible water gas shift (WGS) reaction gets initiated until the exact ratio of the concentration of these gases reaches a particular equilibrium constant KWGS that is a function of temperature. The WGS reaction and its equilibrium constant can be written as:WGS Reaction: CO+H2O<=>CO2+H2 ΔH=−40.6 kJ/mol  (1)WGS equilibrium constant:
                              K          WGS                =                                                            [                                  CO                  2                                ]                            ⁡                              [                                  H                  2                                ]                                                                    [                CO                ]                            ⁡                              [                                                      H                    2                                    ⁢                  O                                ]                                              =                      812.9            -                                                            6.628                  ⁢                  e                                +                5                            T                        +                                                            1.001                  ⁢                  e                                +                8                            T                                                          (        2        )            where T is in ° C. From equation (2), it can be observed that KWGS reduces with increasing temperature. This means that processes aimed at converting coal-derived gas to hydrogen at high temperatures are thermodynamically restricted. While catalysts aid in achieving this equilibrium, they cannot alter the value of K to provide a higher hydrogen yield. An effective technique to shift the reaction to the right for enhanced hydrogen generation has been to remove hydrogen from the reaction mixture. This premise has lead to the development of hydrogen separation membranes. However, membranes cannot completely remove hydrogen from the mixture. Any remaining hydrogen would dilute CO2 after its utilization in either a fuel cell or gas turbine.
Another option for driving the WGS reaction forward is to remove CO2 from the reaction mixture by reacting it with CaO. The carbonation reaction can be written as:Carbonation Reaction CaO+CO2→CaCO3(ΔH=−183 kJ/mol)  (3)Under the appropriate reaction temperature, CO2 concentration can be lowered down to ppm levels by reaction (3), thereby enabling the maximum production of hydrogen from carbon via reaction (1). By conducting the reaction such that CO is the limiting reactant, we can ensure complete utilization of the fuel as well. Besides these advantages, CO2 is simultaneously removed from the gas mixture in the form of CaCO3, thereby improving the purity of the hydrogen stream (the other contaminant being only water). The spent sorbent can then be calcined separately to yield pure CO2 stream, which is then amenable for compression and liquefaction before its transportation to sequestration sites. Calcination reaction, reverse of the carbonation reaction can be written as:Calcination Reaction CaCO3→CaO+CO2 (ΔH=+183 kJ/mol)  (4)The resulting CaO sorbent is recycled to capture CO2 in the next cycle. This cyclical CCR process can be continued so long as the sorbent provides a satisfactory CO2 capture.
To obtain high purity H2, the WGS reaction is generally carried out in two stages for: (1) high temperature shift (250-500° C.) using iron catalysts and (2) low temperature shift (210-270° C.) using copper-based catalysts (Gerhartz, 1993; Bohlbro, 1969). Copper based catalysts are extremely intolerant to small quantities of sulfur (<0.1 ppm) and hence the fuel gases need to be desulfurized upstream of the WGS reactor. Besides, to achieve satisfactory carbon monoxide conversion a considerable quantity of high-pressure steam is required. For example, to lower the CO content of the typical fuel gas from 45% (inlet) to 3% (outlet) a total steam addition of 1.18 kg/m3 of the gas is required, at a total pressure of 60 bar and 410° C. (Gerhartz, 1993). The steam to CO ratio at 550° C. can be as high as 50 during a single-stage operation or 7.5 for a more expensive dual-stage process to obtain 99.5% pure H2 (David, 1980). This is necessary due to the equilibrium limitation inherent in the WGS reaction. From the point of view of H2 production, even though higher temperatures lead to improved kinetics, WGS has poor equilibrium conditions at the higher temperatures. However, the continuous removal of the carbon dioxide product from the reaction chamber will incessantly drive the equilibrium-limited water-gas shift reaction forward. This will ensure a high yield and purity of hydrogen with near stoichiometric amounts of steam needed for the reaction. Besides, the reaction can now be carried out at higher temperatures leading to superior kinetics in the forward direction. Thus the major equilibrium related drawback in this process could be overcome. The continuous CO2 removal can be brought about by the carbonation reaction of a metal oxide to give the corresponding metal carbonate. We have identified a high reactivity, mesoporous calcium oxide as the potential sorbent for the in-situ CO2 capture given by eqn. 3.
The success of this process would effectively bridge coal gasification to fuel cell usage and chemical synthesis. Other side benefits of this process involve the potential for removal of sulfur and heavy metals such as arsenic and selenium from the fuel gas stream.
Recently, Harrison and co-workers reported a single-step sorption-enhanced process to produce hydrogen from methane (Balasubramanian et al., 1999; Lopez Ortiz and Harrison, 2001). They used the traditional concept of SMR with WGS using Ni-based catalyst to produce hydrogen, coupled with this novel scheme of in-situ continuous CO2 capture using a calcium-based dolomite sorbent. They obtained high hydrogen yields with 97% purity (dry basis).
However, they reported a low “calcium” conversion in the sorbent of about 50% at the beginning of the breakthrough to about 83% at the end of the test. These conversion calculations are based on only the calcium portion of their dolomite sorbent. Their total sorbent conversion will be much lower than these values as dolomite does not entirely contain calcium based material. In fact, dolomite comprises of nearly 50 wt. % calcium, which participates in the reaction to some extent, and the remaining portion of the sorbent (mainly magnesium oxide) stays unreacted. Further, they attribute the incomplete conversions of the calcium material to the concept of pore filling and pluggage at the pore-mouths of these sorbent particles by CaCO3 product layer, preventing the access of CO2 in the gas to unreacted CaO surface at the pore interiors.
Harrison and co-workers regenerated the dolomite sorbent in streams of N2, 4% O2 in N2 and pure CO2. They had to use high regeneration temperatures of 800-950° C., especially while using pure CO2. Exposure of the reforming catalyst to an oxidizing atmosphere (viz. O2/N2 or CO2) while regenerating the sorbent used to oxidize the Ni catalysts to NiO. Hence, the catalyst had to be reduced back to Ni before every cycle or the sorbent-catalyst mixture had to be separated after every run so that only the sorbent is subjected to the regeneration conditions. Further, the temperature of operation can be lowered by regeneration in a pure N2 stream. However, it would not solve the problem of CO2 separation due to the formation of a CO2/N2 gas mixture. Calcination in a pure CO2 stream will result in higher operating temperatures due to the thermodynamic limitations of the calcination reaction in presence of the CO2 product. Higher temperatures and the presence of CO2 during calcination would cause the sorbent to sinter. This is in agreement with the results of multiple carbonation-calcination cycle tests for dolomite by Harrison and co-workers (Lopez Ortiz and Harrison, 2001) in pure CO2 stream (800-950° C.). They observed a decrease in “calcium” conversion from 83% in the 1st cycle to about 69% in the 10th cycle itself. However, a mesoporous high suface area calcium based sorbent (precipitated calcium carbonate, PCC) developed at OSU has undergone 100 cycle experiments. The PCC sorbent has shown 85% conversion in the 1st cycle 66.7% in the 10th cycle and 45.5% in the 100th cycle towards carbonation. These experiments were carried out in a TGA at 700° C. in a 10% CO2 stream in the carbonation cycle and 100% N2 gas in the calcination cycle, with 30 minute residence times for each cycle. Therefore this project aims testing this PCC based sorbent towards further enhancing the WGSR and overcoming some of the problems faced by Harrison and co-workers.